This category has the following 4 subcategories, out of 4 total.
Pages in category "Fluorine"
The following 20 pages are in this category, out of 20 total. This list may not reflect recent changes (learn more).
This category has the following 4 subcategories, out of 4 total.
The following 20 pages are in this category, out of 20 total. This list may not reflect recent changes (learn more).
1. Fluorine – Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a highly toxic pale yellow diatomic gas at standard conditions, as the most electronegative element, it is extremely reactive, almost all other elements, including some noble gases, form compounds with fluorine. Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance, proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, and several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, the rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite which plays a key role in aluminium refining. Organic fluorides have very high chemical and thermal stability, their uses are as refrigerants, electrical insulation and cookware. Pharmaceuticals such as atorvastatin and fluoxetine also contain fluorine, and the fluoride ion inhibits dental cavities, global fluorochemical sales amount to more than US$15 billion a year. Fluorocarbon gases are generally greenhouse gases with global-warming potentials 100 to 20,000 times that of carbon dioxide, organofluorine compounds persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no metabolic role in mammals, a few plants synthesize organofluorine poisons that deter herbivores. Fluorine atoms have nine electrons, one fewer than neon, and electron configuration 1s22s22p5, the outer electrons are ineffective at nuclear shielding, and experience a high effective nuclear charge of 9 −2 =7, this affects the atoms physical properties. Fluorines first ionization energy is third-highest among all elements, behind helium and neon and it also has a high electron affinity, second only to chlorine, and tends to capture an electron to become isoelectronic with the noble gas neon, it has the highest electronegativity of any element. Fluorine atoms have a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen, conversely, bonds to other atoms are very strong because of fluorines high electronegativity. Unreactive substances like powdered steel, glass fragments, and asbestos fibers react quickly with cold fluorine gas, wood, reactions of elemental fluorine with metals require varying conditions. Some solid nonmetals react vigorously in liquid air temperature fluorine, hydrogen sulfide and sulfur dioxide combine readily with fluorine, the latter sometimes explosively, sulfuric acid exhibits much less activity, requiring elevated temperatures. Hydrogen, like some of the metals, reacts explosively with fluorine. Carbon, as black, reacts at room temperature to yield fluoromethane. Graphite combines with fluorine above 400 °C to produce non-stoichiometric carbon monofluoride, higher temperatures generate gaseous fluorocarbons, heavier halogens react readily with fluorine as does the noble gas radon, of the other noble gases, only xenon and krypton react, and only under special conditions. At room temperature, fluorine is a gas of diatomic molecules and it has a characteristic pungent odor detectable at 20 ppb
2. Henri Moissan – Ferdinand Frederick Henri Moissan was a French chemist who won the 1906 Nobel Prize in Chemistry for his work in isolating fluorine from its compounds. Moissan was one of the members of the International Atomic Weights Committee. Moissan was born in Paris on 28 September 1852, the son of an officer of the eastern railway company, Francis Ferdinand Moissan. In 1864 they moved to Meaux, where he attended the local school, in 1870 he left the school without the grade universitaire necessary to attend the university. He began working for a chemist in Paris, where he was able to save a person poisoned with arsenic and he decided to study chemistry and began first at the laboratory of Edmond Frémy and later at that of Pierre Paul Dehérain. Dehérain persuaded him to pursue an academic career and he passed the baccalauréat, which was necessary to study at university, in 1874 after an earlier failed attempt. During his time in Paris he became a friend of the chemist Alexandre Léon Étard and he published his first scientific paper, about carbon dioxide and oxygen metabolism in plants, with Dehérain in 1874. After Moissan received his Ph. D. in 1880, his friend Landrine offered him a position at an analytic laboratory and his marriage, to Léonie Lugan, took place in 1882. During the 1880s, Moissan focused on chemistry and especially the production of fluorine itself. He had no laboratory of his own, but used several laboratories, the electrolysis of hydrogen fluoride yielded fluorine on 26 June 1886. After resolving the problem and demonstrating the production of several times. In subsequent years, until 1891, he focused on the study of fluorine chemistry and he discovered numerous fluorine compounds, such as SF6 in 1901. His research in the production of boron and artificial diamonds and the development of a heated oven capable of reaching 3500°C using 2200 amperes at 80 volts followed by 1900. His newly developed arc furnace led to the production of borides and carbides of numerous elements, the existence of the element fluorine had been well known for many years, but all attempts to isolate it had failed, and some experimenters had died in the attempt. Moissan eventually succeeded in preparing fluorine in 1886 by the electrolysis of a solution of potassium hydrogen difluoride in liquid hydrogen fluoride, the mixture was needed because hydrogen fluoride is a nonconductor. The device was built with platinum/iridium electrodes in a platinum holder, the result was the complete isolation of the hydrogen produced at the negative electrode from the fluorine produced at the positive one. This is essentially still the way fluorine is produced today, for this achievement, he was awarded the Nobel Prize in 1906. Late in his life, the government of France named him a Commandeur de la Legion dhonneur, in 1893, Moissan began studying fragments of a meteorite found in Meteor Crater near Diablo Canyon in Arizona
3. History of fluorine – Fluorine is a relatively new element in human applications. In ancient times, only uses of fluorine-containing minerals existed. The industrial use of fluorite, fluorines source mineral, was first described by early scientist Georgius Agricola in the 16th century, the name fluorite derives from Agricolas invented Latin terminology. In the late 18th century, hydrofluoric acid was discovered, by the early 19th century, it was recognized that fluorine was a bound element within compounds, similar to chlorine. Fluorite was determined to be calcium fluoride, because of fluorines tight bonding as well as the toxicity of hydrogen fluoride, the element resisted many attempts to isolate it. In 1886, French chemist Henri Moissan, later a Nobel Prize winner, succeeded in making elemental fluorine by electrolyzing a mixture of potassium fluoride, large-scale production and use of fluorine began during World War 2 as part of the Manhattan Project. Earlier in the century, the main fluorochemicals were commercialized by the DuPont company, refrigerant gases, some instances of ancient use of fluorite, main source mineral of fluorine, for ornamental use carvings exist. However, archeological finds are rare, perhaps in part because of the stones softness, two Roman cups made of Persian fluorite have been discovered and are currently exhibited at the British museum. Pliny the Elder described a stone from Persia used in cups that may have been fluorite. Fluorite carvings from about 1000 AD have been discovered in the Americas in Indian burial grounds, the word fluorine derives from the Latin stem of the main source mineral, fluorite, which was first mentioned in 1529 by Georgius Agricola, the father of mineralogy. He described fluorite as an additive that helps melt ores. Fluorite stones were called schone flusse in the German of the time, Agricola, writing in Latin but describing 16th century industry, invented several hundred new Latin terms. For the schone flusse stones, he used the Latin noun fluores, fluxes, after Agricola, the name for the mineral evolved to fluorspar and then to fluorite. Fluorite mineral was described in the writings of alchemist Basilius Valentinus. However, it is alleged that Valentinus was a hoax as his writings were not known until about 1600, some sources claim that the first production of hydrofluoric acid was by Heinrich Schwanhard, a German glass cutter, in 1670. A peer-reviewed study of Schwanhards writings, though, showed no specific mention of fluorite and it was hypothesized that this was probably nitric acid or aqua regia, both capable of etching soft glass. Andreas Sigismund Marggraf made the first definite preparation of hydrofluoric acid in 1764 when he heated fluorite with sulfuric acid in glass, in 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction. Scheele recognized the product of the reaction as an acid, which he called fluss-spats-syran, in English, in 1810, French physicist André-Marie Ampère suggested that hydrofluoric acid was a compound of hydrogen with an unknown element, analogous to chlorine
4. Biological aspects of fluorine – Among the most reactive of the elements, it has proven valuable in many potent industrial compounds that are quite dangerous to living organisms, such as the weak acid hydrogen fluoride. Fluorine is a component of so-called 1080 poison, a mammal-killer banned in much of the world but still used to control populations of Australian foxes, because carbon-fluorine bonds are difficult to form, they are seldom found in nature. A few species of plants and bacteria found in the tropics make fluorine-containing poisons to deter predators from eating them, when applied topically in dental products, the fluoride ion chemically binds to surface tooth enamel, making it marginally more acid-resistant. Although politically controversial, fluoridation of water supplies has shown consistent benefits to dental hygiene. Manmade fluorinated compounds have also played a role in several environmental concerns. Chlorofluorocarbons, once major components of numerous commercial products, have proven damaging to Earths ozone layer. Similarly, the stability of many organofluorines has raised the issue of biopersistence, long-lived molecules from waterproofing sprays, for example PFOA and PFOS, are found worldwide in the tissues of wildlife and humans, including newborn children. Fluorine biology is relevant to a number of cutting-edge technologies. PFCs are capable of holding enough oxygen to support human liquid breathing, several works of science fiction have touched on this application, but in the real world, researchers have experimented with PFCs for burned lung care and as blood substitutes. Fluorine in the form of its radioisotope 18F is also at the heart of a medical imaging technique known as positron emission tomography. A PET scan produces three-dimensional colored images of parts of the body that use a lot of sugar, biologically synthesized organofluorines have been found in microorganisms and plants, but not in animals. The most common example is fluoroacetate, with an active poison molecule identical to commercial 1080, in bacteria, the enzyme adenosyl-fluoride synthase, which makes the carbon–fluorine bond, has been isolated. The discovery was touted as possibly leading to biological routes for organofluorine synthesis, fluoride is not considered an essential mineral element for mammals and humans. Small amounts of fluoride may be beneficial for bone strength, since the mid-20th century, it has been discerned from population studies that fluoride reduces tooth decay. Initially, researchers hypothesized that fluoride helped by converting tooth enamel from the more acid-soluble mineral hydroxyapatite to the less acid-soluble mineral fluorapatite, however, more recent studies showed no difference in the frequency of caries amongst teeth that were pre-fluoridated to different degrees. Current thinking is that fluoride prevents cavities primarily by helping teeth that are in the early stages of tooth decay. When teeth begin to decay from the acid of sugar-consuming bacteria, however, teeth have a limited ability to recover calcium if decay is not too far advanced. Fluoride appears to reduce demineralization and increase remineralization, also, there is some evidence that fluoride interferes with the bacteria that consume sugars in the mouth and make tooth-destroying acids
5. Fluorochemical industry – The global market for chemicals from fluorine was about US$16 billion per year as of 2006. The industry was predicted to reach 2.6 million metric tons per year by 2015, the largest market is the United States. Western Europe is the second largest, asia Pacific is the fastest growing region of production. China in particular has experienced significant growth as a market and is becoming a producer of them as well. Fluorite mining was estimated in 2003 to be a $550 million industry, mined fluorite is separated into two main grades, with about equal production of each. Acidspar is at least 97% CaF2, metspar is much lower purity, metspar is used almost exclusively for iron smelting. Acidspar is primarily converted to hydrofluoric acid, the resultant HF is mostly used to produce organofluorides and synthetic cryolite. About 3 kg of metspar grade fluorite, added directly to the batch, are used for every ton of steel made. The fluoride ions from CaF2 lower the temperature and viscosity. The calcium content has a benefit in removing sulfur and phosphorus. Metspar is similarly used in cast iron production and for other iron-containing alloys, fluorite of the acidspar grade is used directly as an additive to ceramics and enamels, glass fibers and clouded glass, and cement, as well as in the outer coating of welding rods. Acidspar is primarily used for making hydrofluoric acid, which is an intermediate for most fluorine-containing compounds. Significant direct uses of HF include pickling of steel, cracking of alkanes in the petrochemical industry, one third of HF is used to make synthetic cryolite and aluminium trifluoride. These compounds are used in the electrolysis of aluminium by the Hall–Héroult process, about 23 kg are required for every metric ton of aluminium. These compounds are used as a flux for glass. Fluorosilicates are the next most significant inorganic fluorides formed from HF, the most common one, that of sodium, is used for water fluoridation, as an intermediate for synthetic cryolite and silicon tetrafluoride, and for treatment of effluents in laundries. MgF2 and, to an extent, other alkaline earth difluorides are specialty optical materials. Magnesium difluoride is used as an antireflection coating for spectacles
6. Origin and occurrence of fluorine – Fluorine is relatively rare in the universe compared to other elements of nearby atomic weight. On earth, fluorine is essentially only in mineral compounds because of its reactivity. The main commercial source, fluorite, is a common mineral, at 400 ppb, fluorine is estimated to be the 24th most common element in the universe. It is comparably rare for a light element, all of the elements from atomic number 6 to atomic number 14 are hundreds or thousands of times more common than fluorine except for 11. One science writer described fluorine as a shack amongst mansions in terms of abundance, fluorine is so rare because it is not a product of the usual nuclear fusion processes in stars. And any created fluorine within stars is rapidly eliminated through strong nuclear fusion reactions—either with hydrogen to oxygen and helium, or with helium to make neon. The presence of fluorine at all—outside of temporary existence in stars—is somewhat of a mystery because of the need to escape these fluorine-destroying reactions. Three theoretical solutions to the mystery exist, In type II supernovae, atoms of neon could be hit by neutrinos during the explosion, in Wolf-Rayet stars, a strong solar wind could blow the fluorine out of the star before hydrogen or helium could destroy it. Finally, in asymptotic giant branch stars, fusion reactions occur in pulses, only the red giant hypothesis has supporting evidence from observations. In space, fluorine commonly combines with hydrogen to form hydrogen fluoride, in addition to HF, monatomic fluorine has been observed in the interstellar medium. Fluorine cations have been seen in planetary nebulae and in stars, fluorine is the thirteenth most common element in Earths crust, comprising between 600 and 700 ppm of the crust by mass. Because of its reactivity, it is only found in compounds. Three minerals exist that are industrially relevant sources of fluorine, fluorite, fluorapatite, fluorite, also called fluorspar, is the main source of commercial fluorine. Fluorite is a colorful mineral associated with hydrothermal deposits and it is common and found worldwide. Canada also exited production in the 1990s, the United Kingdom has declining fluorite mining and has been a net importer since the 1980s. Fluorapatite is mined along with other apatites for its content and is used mostly for production of fertilizers. Most of the Earths fluorine is bound in this mineral, but because the percentage within the mineral is low, only in the United States is there significant recovery. There, the hexafluorosilicates produced as byproducts are used to water fluoridation
7. Phases of fluorine – Fluorine forms diatomic molecules that are gaseous at room temperature with a density about 1.3 times that of air. Though sometimes cited as yellow-green, pure fluorine gas is actually a pale yellow. The element has a pungent characteristic odor that is noticeable in concentrations as low as 20 ppb, fluorine condenses to a bright yellow liquid at −188 °C, which is near the condensation temperatures of oxygen and nitrogen. The solid state of fluorine relies on Van der Waals forces to hold molecules together, consequently, the solid state of fluorine is more similar to that of oxygen or the noble gases than to those of the heavier halogens. Fluorine solidifies at −220 °C into a structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules, at −228 °C fluorine undergoes a solid–solid phase transition into a monoclinic structure called alpha-fluorine. This phase is opaque and hard, with close-packed layers of molecules, the solid state phase change requires more energy than the melting point transition and can be violent, shattering samples and blowing out sample holder windows. Solid fluorine received significant study in the 1920s and 30s, the crystal structure of alpha-fluorine given, which still has some uncertainty, dates to a 1970 paper by Linus Pauling. Jordan, T. H. Streib, W. E. Lipscomb, jordan, T. H. Streib, W. D. Smith, H. W. Lipscomb, W. N. Single-crystal studies of β-F2and of γ-O2. Etters, R. D. Kirin, D. High-pressure behavior of molecular fluorine at low temperatures. Kobashi, K. Klein, M. L. Lattice vibrations of solid α-F2, english, C. A. Venables, J. A. The Structure of the Diatomic Molecular Solids, proceedings of the Royal Society A, Mathematical, Physical and Engineering Sciences
8. Covalent radius of fluorine – The covalent radius of fluorine is a measure of the size of a fluorine atom, it is approximated at about 60 picometres. Since fluorine is a small atom with a large electronegativity. The covalent radius is defined as half the bond lengths between two atoms of the same kind connected with a single bond. By this definition, the covalent radius of F is 71 pm, however, the F-F bond in F2 is abnormally weak and long. Bonds to fluorine have considerable ionic character, a result of its atomic radius. Therefore, the length of F is influenced by its ionic radius, the size of ions in an ionic crystal. The ionic radius of fluoride is much larger than its covalent radius, when F becomes F−, it gains one electron but has the same number of protons, meaning the attraction of the protons to the electrons is weaker, and the radius is larger. The first attempt at trying to find the covalent radius of fluorine was in 1938, Brockway prepared a vapour of F2 molecules by means of the electrolysis of potassium bifluoride in a fluorine generator, which was constructed of Monel metal. Then, the product was passed over potassium fluoride so as to remove any hydrogen fluoride, a sample was collected by evaporating the condensed liquid into a Pyrex flask. Finally, using electron diffraction, it was determined that the length between the two fluorine atoms was about 145 pm. He therefore assumed that the covalent radius of fluorine was half this value and this value, however, is inaccurate due to the large electronegativity and small radius of fluorine atom. In 1941, Schomaker and Stevenson proposed an equation to determine the bond length of an atom based on the differences in electronegativities of the two bonded atoms. DAB = rA + rB – C|xA – xB| This equation predicts a bond length which closer to the experimental value and its major weakness is the use of the covalent radius of fluorine that is known as being too large. In 1960, Linus Pauling proposed an additional effect called back bonding to account for the experimental values compared to the theory. His model predicts that F donates electrons into a vacant atomic orbital in the atom it is bonded to, giving the bonds a certain amount of sigma bond character. In addition, the fluorine atom also receives an amount of pi electron density back from the central atom giving rise to double bond character through π or π back bonding. Thus, this suggests that the observed shortening of the lengths of bonds is due to these double bond characteristics. Reed and Schleyer, who were skeptical of Pauling’s proposition, suggested another model in 1990 and they determined that there was no significant back-bonding, but instead proposed that there is extra pi bonding, which arose from the donation of ligand lone pairs into X-F orbitals